Abstract

When you burn a piece of wood, you start with wood and oxygen, and end up with ashes, carbon dioxide, and water vapor. At no time during the reaction does wood reappear, even momentarily, from the ashes. Most chemical reactions are like this; they move in one direction, from reactants (starting chemicals) to products. In this chemistry science project, you will experiment with a rare and exotic reaction that oscillates. The reaction products appear and disappear for a number of cycles. Because the products are colored, the solution appears alternately blue, then yellow, then clear. The reaction is easy to set up, fun to watch, and opens up lots of ways to explore the nature of chemical reactions. Although it would be helpful if you have had a class in chemistry, you can still do this science project even if you have not.

APA Style

Science Buddies Staff.
(2014, November 13).
Minds of Their Own: A Chemical Reaction that Changes, then Changes Back!.
Retrieved March 31, 2015
from http://www.sciencebuddies.org/science-fair-projects/project_ideas/Chem_p097.shtml

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Introduction

What do these things have in common: rust forming on an iron nail, a mixture of vinegar and baking soda producing carbon dioxide bubbles, and gasoline burning in a car's engine? They are all examples of chemical reactions. In each example, the starting chemicals, or reactants, combine to form the resulting chemicals, or products. Rust (or iron oxide) forms when iron in the metal combines with oxygen in the atmosphere. Carbon dioxide forms when the acetic acid in vinegar reacts with sodium bicarbonate in baking soda. And water, carbon dioxide, and the energy used to make a car move result when gasoline reacts with oxygen. These reactions all move in one direction, from reactants to products. However, the Briggs-Rauscher (BR) reaction is different from these reactions in that it oscillates. To start the chemical reaction, two clear solutions are mixed together.

The resulting clear solution then turns blue, then yellow, and then clear again. Each color is present for about 1–3 seconds (sec). The cycle of color changes repeats until one or more of the chemicals are used up. The reaction was developed by Thomas S. Briggs and Warren C. Rauscher of Galileo High School in San Francisco in 1972.

Don't let the equations below scare you away! Don't worry if you think you'll find the following chemistry explanations and chemical equations a bit intimidating—most professional chemists would agree that the chemistry of the Briggs-Rauscher reaction is complicated! While you will not need to use the following equations to perform the experiment below, the information is included so you can try to learn and understand the reactions taking place. There are a lot of things going on at the molecular level to create the oscillating color changes. However, this science project has the advantage that you can actually see changes in the reaction products (the colored ones, at least) as they form and disappear. The Experimental Procedure, below, focuses on how changing the concentration of one of the chemicals (malonic acid) affects the color changes in the reaction. The Variations at the end suggest other experiments if you want to take the project further.

The following explanation of the chemistry involved in the reaction is based on the University of Leeds chemistry website and Shakhashiri's book (see the Bibliography). In the BR reaction, the evolution of oxygen and carbon dioxide gases and the concentrations of iodine and iodide ions oscillate. Here is the list of names for the chemicals in the reactions that follow:

IO3-: Iodate ion, from sodium iodate

H2O2: Hydrogen peroxide

CH2(COOH)2: Malonic acid

H+: Hydrogen ion

ICH(COOH)2: Iodomalonic acid

O2: Oxygen

H2O: Water

HIO: Hypoiodous acid

I-: Iodide ion

HIO2: Iodous acid

Mn2+: Manganese ion, from manganese sulfate

Mn(OH)2: Manganese dioxide

HOO.: Hydroperoxyl radical

IO2.: Iodide dioxide radical

The mechanism of this reaction can be summarized by the following equations:

Equation 1:

IO3- + 2 H2O2 + CH2(COOH)2 + H+ →
ICH(COOH)2 + 2 O2 + 3 H2O

This transformation can be represented by two component reactions:

Equation 2:

IO3- + 2 H2O2 + H+ →
HIO + 2 O2 + 2 H2O

Equation 3:

HIO + CH2(COOH)2 →
ICH(COOH)2 + H2O

The first of these two reactions can occur by either of two different processes, a radical process and a non-radical process (radicals are atoms, molecules, or ions with unpaired electrons, represented as a dot after the name, as in HOO., the hydroperoxyl radical). These two component reactions "compete" for dominance, and the processes that dominates is determined by the concentration of iodide ions in the solution. When [I-] is low, the reaction proceeds primarily by the radical process; when [I-] is high, the non-radical process is the major process. The second reaction (Equation 3) couples the two processes. The reaction consumes HIO more slowly than HIO is produced by the radical process when that process is predominant, but it uses up HIO more rapidly than it is produced by the non-radical process. Any HIO that does not react by Equation 3 is reduced to I- by hydrogen peroxide as one of the component steps of the non-radical process for reaction 2.

When HIO is produced rapidly by the radical process, the excess forms the iodide ions, which shut off that radical process and start the slower non-radical process. Reaction 3 then consumes the HIO so rapidly that not enough is available to produce the iodide ion necessary to keep the nonradical process going, and the radical process starts again. Each of the processes of reaction 2 produces conditions favorable to the other process; therefore, the reaction oscillates between these two processes.

Let's look at the process in a little more detail. If iodide ions are present in sufficient concentration, the reaction follows the non-radical process, reaction 2. The iodide ions react relatively slowly with iodate ions.

Because reaction 2 is slower than reaction 3 under these conditions, so much HIO is used up by reaction 3 that reaction 6 cannot replenish the I- consumed in reactions 4 and 5; the [I-] keeps diminishing.

Once the concentration of iodide ions falls below a certain level, the non-radical process becomes very slow, and the radical process for reaction 2 takes over. This process involves these five steps:

Equation 7:

IO3- + HIO2 + H+ → 2 IO2· + H2O

Equation 8:

IO2· + Mn2++ H2O → HIO2 + Mn(OH)2

Equation 9:

Mn(OH)2 + H2 O2 → Mn2+ + H2O + HOO·

Equation 10:

2 HOO· → H2O2 + O2

Equation 11:

2 HIO2 → IO3- + HIO + H+

These steps, when combined in the stoichiometry of 2 (Equation 7) + 4 (Equation 8) + 4 (Equation 9) + 2 (Equation 10) +1 (Equation 11), have the overall result given by Equation 2. A significant feature of this process is that, taken together, the first two steps, Equation 7 and Equation 8, are autocatalytic—they produce 2 HIO2 for each one consumed. Therefore, the rate of these steps increases as they occur. Because this radical process is autocatalytic, it causes a rapid increase in the concentration of HIO, which is produced by the disproportionation of HIO2 (Equation 11). This process does not rapidly consume all the iodate in the solution, because the last step is second order in the catalytic species. Thus, as its concentration increases because of the autocatalytic nature of the early steps, HIO2 is ever more rapidly consumed in this last step, and the sequence of the reactions quickly reaches a steady state.

Equations 8 and 9 show the function of the manganese catalyst. The manganese is oxidized in reaction 8 and reduced in reaction 9. Its catalytic effect in the reaction is accounted for through its providing the means for reducing IO2· radicals to HIO2, thereby completing the autocatalytic cycle of equations 7 and 8.

The hypoiodous acid produced by the radical process reacts with malonic acid by reaction 3. However, the radical process is faster than reaction 3, and the excess HIO reacts with hydrogen peroxide by reaction 6 to create I-, which shuts off the radical process and returns the system to the slow nonradical process initiated by reaction 4.

The dramatic color effects arise because reaction 3 does not take place in a single step, but by the sequence of reactions 12 and 13.

Equation 12:

I- + HIO + H+ → I2 + H2O

Equation 13:

I2 + CH2(COOH)2 → ICH(COOH)2 + H+ + I-

The solution turns amber from the I2 produced through reaction 1, when the radical process maintains [HIO] greater than [I-]. The excess HIO is converted to I- through the reaction with H2O2 (Equation 6). The solution suddenly turns dark blue when [I-] becomes greater than [HIO], and the I- can combine with I2 to form a complex with the starch. With [I-] high, reaction 2 switches to the slow nonradical process. The color then fades as reaction 3 consumes iodine faster than it is produced. When the system switches back to the rapid radical process, the cycle repeats.

When the solutions containing the reactants are mixed, IO3- reacts with H2O2 to produce a little HIO2. The HIO2 reacts with IO3- in the first step of the radical process (Equation 7). The autocatalytic radical process follows, rapidly increasing the concentration of HIO. The HIO is reduced to I- in a reaction with H2O2 (Equation 6). The large amount of HIO reacts with I-, producing I2 (Equation 12). The I2 reacts slowly with malonic acid, but the concentration of HIO, I2- all increase, because reaction 2 is faster than reaction 3. As [I-] increases, the rate of its reaction with HIO2 (Equation 5) surpasses that of the autocatalytic sequence of reactions 7 and 8. The radical process is then shut off, and the accumulation of reduced iodine is consumed by reaction 3 operating through the sequence of reactions 12 and 13. Eventually, [I-] is reduced to such a low value that reactions 7 and 8 become faster than reaction 5, and the radical process takes over again. This oscillating sequence repeats until the malonic acid or IO3- is depleted.

For this science project, you will buy a kit that contains all of the chemicals you need. Wear gloves and safety goggles when working with chemicals. The focus of the procedure is to investigate how changing the amount of malonic acid affects the color changes produced in the reaction.

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Materials and Equipment

The Fascinating Oscillating Reaction Kit, Educational Innovations, Inc.; available from www.teachersource.com. The kit comes in two sizes:

Item CK-475 has one set of chemicals and provides sufficient chemicals for three trials of the basic procedure.

Item CK-480 is a larger classroom kit that has more chemicals. Order the larger kit if you want to have chemicals for variations of the basic procedure. The classroom kit is also available through Carolina Biological Supply Company (item # 840443). Note: If you are ordering this kit through Carolina Biological Supply Company, the kit must be ordered by a teacher and shipped to a school or business address, so plan accordingly.

The kit contains the following items:

Malonic acid

Manganese sulfate

Sodium iodate

Sodium thiosulfate

Sulfamic acid

Starch solution

Wooden rack

Pipet

Stirrers

Spoon

Small scoops

Cups

Instructions

Hydrogen peroxide (3%) is required and not supplied in the kit; it is available from Carolina Biological Supply Company (item # 868093) or at any drug store. Note: If you are ordering this chemical through Carolina Biological Supply Company, the chemical must be ordered by a teacher and shipped to a school or business address, so plan accordingly.

Permanent marker

Distilled water; available at grocery or drug stores

Stopwatch

Lab notebook

Adult helper

Optional: Video recorder to record the color changes in the chemical reaction, allowing you to obtain more accurate times when you view it later

Graph paper

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Experimental Procedure

Arrange the chemicals in the wooden rack to make them easier to work with.

The chemicals you will need are sodium iodate, sulfamic acid, malonic acid, and manganese sulfate. They are in plastic vials.

You will make enough of the solutions so that you can perform the procedure three times with the original solutions. This prevents variation due to measuring out different amounts of chemicals in separate trials.

You will make an excess of each solution (40 mL rather than 30) so that there is enough to run it three times with an equal amount of liquid. If you make just enough for three trials, the volumes available for the third trial might be less than needed.

Making the Original Solutions

Making Solution A

Label one cup A.

Pour 40 mL of distilled water into cup A using the beaker.

Add two scoops of sodium iodate and four level scoops of sulfamic acid to the cup.

Swirl gently until the chemicals are dissolved.

This is solution A.

Making Solution B

Label another cup B.

Add 40 mL of hydrogen peroxide to cup B using the beaker.

Add four rounded scoops of malonic acid to cup B.

Add a small amount of manganese sulfate, about the size of a grain of rice.

Shake the starch solution bottle, then add 10 drops of the starch solution to cup B.

Swirl the solution to mix the chemicals.

This is solution B.

Mixing the Solutions

Label a third cup C. Be ready to observe and time the reaction process as soon as the two solutions are combined.

Pour 10 mL of solution A into cup C using the beaker.

Rinse the beaker out between uses with tap water.

Pour 10 mL of solution B into cup C.

Start the stopwatch immediately.

Swirl the combined solutions in cup C to mix.

Record the times at which the solution turns blue in your lab notebook.

It should turn blue/clear a number of times. The precise number of cycles depends on the starting conditions.

Have your helper write down the time the solution turned blue in the lab notebook while you watch the solution and the stopwatch. Or work out your own way to work together to record the times at which the solution turns blue.

You could also use a video recorder to capture the changes in color if you choose. This will allow you determine the times at which color changes occurred more accurately.

Wait until the reaction stops cycling between colors. When the reaction stops, the solution will be a brown/purple color.

Carefully pour this solution down the sink with plenty of cold running water to wash it down. Be careful with this solution, as it contains iodine and will stain surfaces with which it comes in contact.

Write down additional observations in your lab notebook. What was the sequence of color changes? Describe the colors as accurately as possible. How quickly did the cycles go?

Repeat the steps 1–9 two more times using solutions A and B. Remember, there will be some of solution A and solution B left over, since you made excess to ensure that you could do the procedure three times with equal amounts of liquid.

You could repeat the procedure a fourth time with the remaining solutions, or dispose of them down a sink with cold running water.

Using Fewer Scoops of Malonic Acid in Solution B

Making Solution A

Clean the cup labeled A with tap water.

Pour 40 ml of distilled water into cup A using the beaker.

Add two scoops of sodium iodate and four level scoops of sulfamic acid.

Add a small amount of manganese sulfate, about the size of a grain of rice.

Shake the starch solution bottle, then add 10 drops of the starch solution to cup B.

Swirl the solution to mix the chemicals.

Mixing the Solutions

Mix solutions A and B in cup C. Get ready to observe and time the reaction process as soon as the two solutions are combined.

Clean the cup labeled C with tap water.

Pour 10 mL of solution A into cup C using the beaker.

Pour 10 mL of solution B into cup C.

Start the stopwatch immediately.

Swirl the combined solutions in cup C to mix.

Record the times at which the solution turns blue in your lab notebook.

Repeat steps 1–7 two more times with the remaining solutions A and B.

Carefully dispose of the solutions down a sink with cold running water.

You should now continue to repeat Making Solution A through Mixing the Solutions of this section with the following amounts of malonic acid in step 3 of Making Solution B: two scoops, one scoop, then no scoops. You should end up with data for three trials for each of the different malonic acid amounts.

Analyze Your Results

Average the times at which you observed the color change for each trial with a given amount of malonic acid.

Graph the average times at which the color changed to blue for each amount of malonic acid on the x-axis.

Make five horizontal lines parallel to the x-axis. Label them 0, 1, 2, 3, and 4 to indicate the number of scoops of malonic acid.

Put a dot on each line at the time the color changed to blue.

What happened as you decreased the amount of malonic acid in the reaction?

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Variations

Vary other chemical components and discuss your observations about what happens in each case.

What happens if there is no starch?

The reaction generates gas. How would increasing the pressure above the reaction affect the reaction rate? Devise a way to test your prediction.

Vary the temperature; for example, 0, 10, 20, and 30°C.

Analyze the timing of the reaction using a digital recorder.

The oscillations of the Briggs-Rauscher reaction depend on the presence of free radicals in the solution. Adding antioxidants has been reported to change the oscillations due to their ability to "scavenge" the free radicals. Devise a way to use the Briggs-Rauscher reaction to test a chemical's ability to function as an antioxidant.

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Ask an Expert

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